Chapter 2: Atoms, Molecules & Life What Are Atoms? An atom are the smallest unit of matter. Atoms are composed of Electrons = negatively charged particles. Neutrons = particles with no charge (neutral). Protons = positively charged particles. Protons and neutrons are found in the nucleus of an atom Electrons orbit the nucleus. An atom is neutral, the # of electrons = # protons An example of an atom: Remember that neutrons are neutral, so there can be more of them than protons. 1
Chapter 2: Atoms, Molecules & Life Periodic Table of Elements: Element = Substance that can t be broken down or converted to another, simpler substance by ordinary chemical means. Atomic Number = Number of protons in the nucleus Since # of protons always equals the # of electrons, why don t we use the # of electrons as the atomic number? Atomic Mass = Number of protons & neutrons in nucleus Periodic Table Why is the mass number a decimal? Chapter 2: Atoms, Molecules & Life Isotope = The same element with a different number of neutrons C 14 P 32 U 235 2
Isotopes Isotopes: Same element, different number of neutrons. Different number of neutrons changes the atomic mass, but NOT the atomic number. Atomic number remains 1 for hydrogen and its isotopes Isotopes Some isotopes, but not all, are radioactive. Example: Carbon 14 (C 14 ) is radioactive Example: Hydrogen 2 (not radioactive) and hydrogen 3 (radiocative) Isotopes are useful in research Nuclear experiments involved heavy water (H 2 ) Radiolabelling used to be H 3, but now other isotopes are used. Radiocarbon Dating: Technique for determining the age of materials that contain carbon based on C 14 levels C 12 O 2 C 14 O 2 Half-life ~ 5730 years 1 C 14 for every 1,000,000,000,000 C 12 3
Crucial elements in life Carbon All organic matter has carbon. 18.5% of the human body mass is carbon atoms. Hydrogen All macromolecules have hydrogen as a component. 9.5% of the human body mass is hydrogen atoms. Oxygen All macromolecules have oxygen as a component. 65% of the human body mass is oxygen atoms. Why so much oxygen? Nitrogen All proteins have nitrogen as a component 3.3% of the human body mass is nitrogen atoms (mainly in muscle and other proteins) Other important elements in life Calcium Component of bones. Phosphorus A component of all cells (phospholipids). Potassium An important electrolyte, also keeps cell alive via sodium potassium pump. Sulfur A component of some protein molecules. Sodium Another important electrolyte, sodium ion pumps. Compounds vs. Molecules Compound: A substance made up of different types of atoms. Example: Table salt, NaCl. Molecule: a particle composed of one or more atoms held together by chemical bonds. Example: Table salt, NaCl Also the smallest unit of a compound. Not all molecules are compounds. H 2, O 2, and other diatomic gases are not molecules. Why? 4
How are molecules formed? The number of electrons in the outermost electron shell determine whether an atom is reactive or inert. Carbon: reactive. 4 electrons in outer shell, needs 4 more to fill shell Neon: inert 8 electrons in outer shell, Does not need electrons to fill shell How are molecules formed? How many electrons are needed to fill an electron shell? Depends on which shell. First shell only needs two hydrogen, helium. Second shell needs eight. Third shell needs eight. Inert Atoms = Atoms with their outermost shell either completely full or empty Reactive Atoms = Atoms with their outermost shell only partially filled Hydrogen Carbon Oxygen Helium Neon Argon 5
How are molecules formed? Reactive atoms want to lose or gain electrons to stabilize their outer (valence) shell. (Figure 2.3) Chemical bonds Types of chemical bonds Covalent: the strongest of the three main types. Ionic: weaker than covalent bonds. Hydrogen bonds: the weakest chemical bond of the three main types. Chemical bonds are crucial for chemical reactions Chemical Reaction = making or breaking chemical bonds. Chemical reactions are essential for all life. Covalent bonds: the strongest chemical bond Atoms share electrons in covalent bonds. Sharing of electrons 6
Covalent bonds: the strongest Most biological molecules utilize covalent bonds. Proteins Carbohydrates Lipids DNA Carbon atoms are always linked by covalent bonds. Examples of Covalent Bonds: Share one pair of electrons (H 2 ) Single covalent bond H - H Share two pairs of electrons (O 2 ) Double covalent bond O = O Share three pairs of electrons (N 2 ) Triple covalent bond N N Bonding Patterns (Table 2.3): Carbon (C) Oxygen (O) 7
Non-polar Covalent Bonds: Equal sharing of electrons Polar Covalent Bonds: Unequal sharing of electrons Molecule is electrically neutral, but poles are charged due to differences in nuclear attraction for electrons (electronegativity) Ionic Bonds Attractive force between atoms that have lost or gained electrons Creates ions (negatively or positively charged atoms) Transfer of electrons Example of ionic bonds Sodium chloride (table salt) Sodium wants to lose one electron to stabilize its outer shell Chloride needs to gain one electron to stabilize its outer shell. Sodium Chloride (NaCl - Figure 2.3) 8
Ionic bonds typically occur between atoms that are located on opposite sides of the periodic table. Hydrogen Bonds Attractive force between water molecules due to polar covalent bonds. Electrons are far more attracted to oxygen than to hydrogen atoms. hydrogen atoms have a slight positive charge. Oxygen atoms have a slight negative charge. Makes water a very special molecule. (Figure 2.5) Water = Good Stuff! Life most likely arose in water Living organisms 60-90% water Why is Water so Important to Life? 9
Humans and water Human body is 65% water. The average human can survive several months without food. But you can only survive 3 4 days without water. All life depends on water Search for life on other planets often includes the search for H 2 O Phoenix spacecraft will land in one of Martian ice caps in hopes of finding water and microbial life. Landed May, 2008, found ice on July 31 st, 2008. Why is Water so Important to Life? Importance of Water: 1) Water is an excellent solvent: liquid capable of dissolving other substances in itself Dissolving Ionic Bonds: (Salt) Polar nature of water (Figure 2.6) Solution = Fluid containing dissolved substances 10
Dissolving Polar Covalent Bonds: (e.g. glucose) Hydrophilic Molecules: Water-loving + - Molecules electrically attracted to water Ions, polar molecules Hydrophobic Molecules: Water-fearing Molecules electrically neutral (fats / oils) Molecules tend to clump together in water (Figure 2.7) Hydrophilic vs hydrophobic Hydrophilic Hydrophobic Importance of Water: 2) Water molecules tend to stick together (cohesion): Surface Tension: Tendency for a water surface to resist breaking Adhesion: Tendency for water to stick to walls of surfaces Flow against gravity Walk on water 11
Importance of Water: 3) Water Can Form H + and OH - Ions (ionization): H 2 O H + + OH - Pure water contains equal amounts of H + and OH - HCl Acidic Basic NaOH H + = OH - H + > OH - Water H + = OH - H + < Water OH - Acids & bases disrupt the equilibrium. Acid A substance that increases the [H + ] in a solution Base A substance that decreases the [H + ] in a solution The ph of a solution describes its degree of acidity: (Figure 2.9) 12
Importance of ph to living systems Changes in ph disrupt life chemistry Example: Blood usual ph = 7.4 ph 7.0, 7.8 lethal So how does blood maintain a healthy ph? Buffers Buffers - solutes that act to resist changes to the ph of a solution when H + or OH - is added. Biological fluids use buffers to help maintain correct ph. Buffers maintain a solution at relatively constant ph: Stable ph essential for normal function Buffers either accept or release a H + in response to changes in ph Example: Bicarbonate ion (HCO 3- ) HCO 3 - Bicarbonate ion + Too acidic? H + Hydrogen ion H 2 CO 3 Carbonic Acid Too basic? H 2 CO 3 Carbonic Acid + OH - Hydroxide ion HCO 3 - + Bicarbonate ion H 2 0 Water 13
Importance of Water: 4) Water Moderates Temperature Changes: Background: Temperature = Speed of molecules Slow molecules = Cool temperatures Fast molecules = Warm temperatures A) Water Heats Slowly Energy first initiates breaking of hydrogen bonds Specific Heat = energy needed to heat 1 gram of a substance 1 C Specific Heat Water = 1 cal Specific Heat Alcohol = 0.6 cal Specific Heat Granite = 0.02 cal Importance of Water: 4) Water Moderates Temperature Changes: A) Water heats slowly B) Water is an effective coolant Heat of Vaporization: Heat needed to convert liquid water to water vapor 529 calories/gram (very high!) By evaporating 1 g of water, 539 grams of human body cools 1 C C) Water freezes slowly Moderates the effects of low temperatures D) Water forms ice (less dense than fluid water): Acts as an insulator for life below Why does ice weigh less than liquid water? Water is the most dense (0.9998 g/cm 3 ) at 4 0 C. Below 4 degrees, the water molecules form more hydrogen bonds, forming lattice like crystals. This characteristic is very important for aquatic organisms. Why?? 14
Because ice floats! Any questions? Next Class We will be starting Chapter 3: Biological molecules Remember there are games to use as study tools Especially useful for learning terminology. 15